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Uploaded byMohmmad Amil Rahman
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Ph, buffer solution ph indicator and ph

pH, pH measurement, and buffer solutions can be summarized as follows: 1) pH is a measure of how acidic or basic a solution is, measured on a scale from 0-14. It is determined by the concentration of hydrogen ions [H+] in the solution. 2) pH can be measured using pH strips, pH indicators, or a pH meter. Common pH indicators change color depending on the pH range and can indicate if a solution is acidic or basic. 3) Buffer solutions help maintain pH when small amounts of acid or base are added. They use a weak acid or base combined with its salt. The Henderson-Hasselbalch equation can be used to calculate the pH of buffer solutions

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pH, pH measurement, pH indicator and Buffer solution Mohmmad Amil Rahman S.R. Department of Biochemistry Dr. R.P.G.M.College Kangra at Tanda(H.P.)
pH(Potential of Hydrogen) • It is the negative log of the hydrogen ion concentration. pH = -log [H+] • pH is a unit of measure which describes the degree of • acidity or alkalinity (basic) of a solution. • It is measured on a scale of 0 to 14. • Low pH values correspond to high concentrations of H+ and high pH values correspond to low concentrations of H+.
pH VALUE • The pH value of a substance is directly related to the ratio of the hydrogen ion and hydroxyl ion concentrations. • If the H+ concentration is higher than OH- the material is acidic. • If the OH- concentration is higher than H+ the material is basic. • 7 is neutral, < is acidic, >7 is basic
THE pH SCALE • The pH scale corresponds to the concentration of hydrogen ions. • For example pure water H+ ion concentration is 1 x 10^-7 M, therefore the pH would then be 7.
• Acid:-Any compound which forms H⁺ ions in solution (proton donors) eg: Carbonic acid releases H⁺ ions • Base:-Any compound which combines with H⁺ ions in solution (proton acceptors) eg: Bicarbonate(HCO3⁻) accepts H+ ions
MEASUREMENT OF PH • The pH can be measured by: • ph strips • ph indicators • ph meter
• pH Strips • Less precise • Each indicator is only good for a small pH range (1-2 pH units) • But may be good for field usage, or measuring small volumes, or dealing with noxious samples.
• Litmus dye is a natural pH indicator. Litmus paper is a small strip of paper that has dyes in it. When it touches a base or a dye then the litmus paper changes colors depending on , either it is a base or it is an acid. If a litmus paper is exposed to an acid it either changes its colour to pink or red. If a strip of litmus paper touches a base then the paper changes to a shade of blue. Figure : litmus test
SOME IMPORTANT INDICATORS USED IN A CLINICAL BIOCHEMISTRY LABORATORY ARE LISTED BELOW: sr,. No. INDICATOR Ph range Colour in acidic ph Colour in basic ph 1 Phenophthalein 9.3-10.5 colourless pink 2 Methyl orange 3.1-4.6 red yellow 3 Bromophenol blue 3.0-4.6 yellow blue 4 Methyl red 4.4-6.2 Red yellow 5 Phenol red 6.8 – 8.4 yellow red 6 Litmus 4.5-8.3 red Blue
pH meter Electrode measures H+ concentration Must standardize (calibrate) before using. Actually measuring a voltage – a charge differential – between a control solution and the external fluid.
• Types of pH meter 1. Manual pH meter 2. Digital pH meter Figure : manual pH meter
Figure : Digital pH meter
• Types of probe 1. Glass electrode 2. Reference electrode 3. Combination gel electrode Figure : Reference electrode Figure :Glasselectrode
• Calibration •Calibration should be performed with at least two standard buffer solutions that span the range of pH values to be measured. For general purposes buffers at pH 4 and pH 10 are acceptable. The pH meter has one control (calibrate) to set the meter reading equal to the value of the first standard buffer and a second control (slope) which is used to adjust the meter reading to the value of the second buffer. A third control allows the temperature to be set.
19 Applications of pH meter For the diagnosis of various disorders in human body Agriculture Brewing Pharmaceuticals
Buffers • “Buffers are compounds or mixtures of compounds that by their presence in the solution resist changes in the pH upon the addition of small quantities of acid or alkali.” • Types of buffers: • Generally buffers are of two types; a) Acidic buffers b) Basic buffers
Necessity of a buffer system  Sometimes it is necessary that a solution of a definite pH be prepared and stored. The preservation of such a solution is even more difficult than its preparation. If solution comes in contact with air, it will absorb CO2 and becomes acidic. On the other hand, if solution is stored in a glass bottle, alkaline impurities from the glass may alter its pH.  Due to these reasons, chemical solutions are buffered as the buffer solutions are capable of maintaining pH at some fairly constant value when even small amounts of acid or base are added.
• Acidic Buffers: • An acidic buffer is a combination of weak acid and its salt with a strong base. i.e. Weak acid & salt with strong base (conjugate base). • EXAMPLES: • CH3COOH / CH3COONa • H2CO3 / NaHCO3 • H3PO4 / NaH2PO4 • HCOOH / HCOONa
• Basic Buffers: • A basic buffer is a combination of weak base and its salt with a strong acid. i.e. Weak base & salt with strong acid (conjugate acid). • EXAMPLES: • NH4OH / NH4Cl • NH3 / NH4Cl • NH3 / (NH4)2CO3
• Phosphate Buffers (Double salt buffers): • Besides the two general types of buffers (i.e. acidic & basic), a third appears to exist. This is buffer system composed of two salts: • Monobasic potassium phosphate (KH2PO4) • Dibasic potassium phosphate (K2HPO4).
Mechanism of Action of acidic buffers: • Consider a buffer system of CH3COOH (Weak electrolyte) and CH3COONa (Strong electrolyte). There will be a large concentration of Na+ ions, CH3COONa – ions, and un- dissociated CH3COOH molecules. • When an acid is added If a strong acid (HCl) is added in CH3COOH / CH3COONa buffer, the changes that will occur may be represented as:
• The hydrogen ions yielded by the Hcl are quickly removed as unionized acetic acid, and the hydrogen ion concentration is therefore only slightly affected (because acetic acid produced is very weak as compared to Hcl added).
• When a base is added • If a strong base (NaoH) is added in CH3COOH / CH3COONa buffer, the changes that will occur may be represented as: • The hydroxyl ions yielded by the NaOH are therefore removed as water. The supply of hydrogen ions needed for this purpose being constantly provided by the dissociation of acetic acid.
Mechanism of Action of basic buffers: • Consider a buffer system of NH4OH (Weak electrolyte) and NH4Cl (Strong electrolyte). There will be a large concentration of NH4+ ions, Cl – ions, and un- dissociated NH4OH molecules. • When an acid is added • If a strong acid (HCl) is added in NH4OH / NH4Cl buffer, the changes that will occur may be represented as:
• When a base is added • If a strong base (NaOH) is added in NH4OH / NH4Cl buffer, the changes that will occur may be represented as: • The hydroxyl ions yielded by the NaOH are therefore quickly removed as unionized ammonium hydroxide and the pH of solution is only slightly affected.
• Mechanism of Action of phosphate buffers: • In KH2PO4 / K2HPO4 buffer system, H2PO4 serves as weak acid and HPO4 erves as conjugate base. • When hydronium ions are added, then • When hydroxyl ions are added to this buffer, the following reaction takes place;
BUFFER EQUATION (Henderson – Hasselbalch equation)
• For Acid Buffers: • The pH of acid buffer can be calculated from the dissociation constant, Ka of the weak acid and the concentrations of the acid and salt used. • The dissociation expression of the weak acid can be represented as: • HA↔ H+ + A- Ka = [H+] [A-] / [HA] Or [H+]= Ka [HA] / [A-] ------------- (1)
• A weak acid is only slightly dissociated, and its dissociation is further depressed by the addition of the salt (XA) which provides ion (common ion effect) as a result the equilibrium concentration of the unionized acid is nearly equal to the initial concentration of the acid. • The equilibrium concentration of A- is assumed to be equal to the initial concentration of the salt added since it is completely dissociated. Therefore, in above equation (1), we represent concentration of A- by salt concentration.
• [H+] = Ka. [Acid] / [Salt] --------- (2) • Taking log on both sides, we get: • log[H+]= logKa + log [Acid] / [Salt] • multiplying both sides by –ve sign, • -log[H+] = -logKa - log [Acid] / [Salt] • As -log[H+] = pH & -logKa = pka • pH = pka - log[Acid] / [Salt] OR • pH = pka + log[Salt] / [Acid] ---------- (3) • Eq. (3) is called as Henderson – Hasselbalch equation. It helps in calculating the pH value of buffer solution, • if the concentrations of acid as well as that of the salt are known.
• For Basic Buffers: • Buffer equation for basic buffer can be calculated in same way as that for acidic buffers. • Consider a basic buffer composed of a mixture of weak base (BOH) and its salt (BA). The dissociation constant for base can be written as, • BOH ↔ B+ + OH- Kb = [B+] [OH-] / [BOH] OR [OH-] = Kb [BOH] / [B+] ------------- (1)
• A weak base is only slightly dissociated, and its dissociation is further depressed by the addition of the salt (BA) which provides B+ ion (common ion effect) as a result the equilibrium concentration of the unionized base is nearly equal to the initial concentration of the base. • The equilibrium concentration of B+ is assumed to be equal to the initial concentration of the salt added since it is completely dissociated. Therefore, in above equation (1), we represent concentration of B+ by salt concentration.
[OH-] = Kb. [Base] / [Salt] --------- (2) Taking log on both sides, we get: log[OH-] = logKb + log [Base] / [Salt] multiplying both sides by –ve sign, -log[OH-] = -logKb - log [Base] / [Salt] As -log[OH-] = pOH & -logKb = pkb pOH = pkb – log [Base] / [Salt] Or pOH =pkb + log[Salt] / [Base] ---------- (3)
Significance of Henderson – Hasselbalch equation: • By this equation, the pH of a buffer solution can be calculated from the initial concentrations of the weak acid and the salt provided when ka is given. • However, the Henderson equation for a basic buffer will give pOH, and so pH can be calculated as; pkw = pH + pOH (OR) pH = pkw – pOH pH = 14 – POH • Also, the dissociation constant of a weak acid (pka) or a weak base (pkb) can be calculated by measuring the pH of a buffer solution containing equimolar concentrations of the acid (or base) and the salt.
• EXAMPLE: • Calculate the pH of a buffer solution containing 0.03 moles/litre of acetic acid and 0.1 moles/litre of sodium acetate. Pka for CH3COOH is 4.57. • Solution: • Conc. Of acid = 0.03 M • Conc. Of salt = 0.1 M • So, • pH = pka + log [salt] / [acid] • =4.57 + log 0.1 / 0.03 • = 4.57 + 0.52 • = 5.09 • Result →The pH of the buffer solution containing 0.03 M of acetic acid and 0.1 M of sodium acetate is 5.09.
• Example • Calculate the pH of a buffer solution containing 0.25 moles / litre of formic acid (HCOOH) and 0.10 moles / litre of sodium formate (HCOONa). Ka for formic acid is 1.8 × 10 -4. Solution: Conc. Of acid = 0.25 M Conc. Of salt = 0.10 M Ka = 1.8 × 10 -4 & pka is pka = -log ka = -log 1.8 × 10 -4 = - (log 1.8 × 10 -4) = - (log 1.8 + log 10 -4) = - [0.25 + (-4)] = - (- 3.75) = 3.75 So, pH = pka + log [salt] / [acid] = 3.75 + log 0.10 / 0.25 = 3.75 – 0.397 = 3.34 Result → The pH of a buffer solution containing 0.25 M of formic acid and 0.10 M of sodium formate is 3.34.
BUFFER CAPACITY • The buffer capacity of a buffer solution is “a measure of its magnitude of its resistance to change in the pH on an addition of an acid or a base.” • Buffer capacity is also referred as buffer index, buffer value, buffer efficiency or buffer coefficient. • The buffer capacity represented by ‘β’ may also be defined as: • “The ratio of the increment (amount added) of strong acid or base to the small change in pH (ΔpH) brought about by this addition”. • β = ΔA or ΔB / ΔpH • Where, ΔA or ΔB represents the small increment (in gram equivalents / litre of strong acid or base added) to the buffer to bring about a pH change of ΔpH. • According to the above equation, a solution has a buffer capacity of 1 when one litre of it requires one gram equivalent of a strong acid or base to change the pH by one unit. So, smaller the pH change in a solution upon the addition of an acid or base, greater is the buffer capacity and vice versa.
• Example • Prepare a buffer solution of pH 5 from acetic acid CH3COOH & CH3COONa. pka of CH3COOH is 4.7. • Molar concentration of acid required = 1M, Molar concentration of base required = x M = ? So, by putting above information in equation, we get: pH = pka + log[Salt] / [Acid] 5 = 4.7 + log [x] / [1] 5 – 4.7 = log x – log 1 as log 1 = 0, 0.3 = log x x = log- 0.3 x =2 • Result→ in order to prepare buffer solution of pH 5, acetic acid CH3COOH & sodium acetate CH3COO must be mixed in a ratio of 1M : 2M (log- means anti-log)
APPLICATIONSOF BUFFERS: BUFFER SYSTEM IN BODY FLUIDS